Chapter 2 Electrochemistry of the Chalcogens€¦ · Chapter 2 Electrochemistry of ... M. Bouroushian, Electrochemistry of Metal Chalcogenides, Monographs 57 in Electrochemistry,

Chapter 2Electrochemistry of the Chalcogens

2.1 General References

Because of their multiple oxidation states, the chalcogens, particularly sulfur, canengage in numerous redox couples participating in acid–base, oxidation–reduction,precipitation, and complexation equilibria.

In the anion electrochemical series, sulfur, being the less noble element com-pared to its heavier congeners, occupies an intermediate position between iodineand selenium [(+)F, Cl, Br, I, S, Se, Te(–)]. Selenium, regarded as a metalloid, isa relatively noble element. Tellurium is rather an amphoteric element: it can enterinto solution in the form of both cations and anions. Regarded as a metal, i.e., withrespect to its cations, tellurium occupies a position between copper and mercury.Regarded as a metalloid, i.e., with respect to its anions, it is located on the extremeright of the above series.

A comprehensive survey of the classical electrochemical facts for sulfur, sele-nium, and tellurium, as documented until about 1970, can be found in the reviewsof Zhdanov [1], wherefrom we cite the lists of standard and formal potentials foraqueous solutions, in Tables 2.1, 2.2, and 2.3, respectively. Many of these potentialshave been calculated thermodynamically since the experimental determinations arefew. The listed data are largely drawn from the monograph by Pourbaix (below), andthe interested reader should validate the measurement conditions for the indicatedpotentials or the scatter in their values (not given here in detail). In these tables, theredox systems are assorted by decreasing formal valency of chalcogen in the oxi-dized state, while at a given valency of the oxidized state they appear in the orderof decreasing valency of chalcogen in the reduced state. Latimer [2] has compileduseful aqueous redox transition potential diagrams (reproduced also in Zhdanov’smonographs) that are convenient as a quick guide in practical problems and forperceiving the oxidation–reduction properties of some chalcogen hydride and oxy-chalcogenide species. Standard potentials of chalcogens in non-aqueous media aregenerally not known, at least in a systematic manner.

A standard approach for the theoretical presentation of electrochemical equilibriais the use of Pourbaix, or potential–pH predominance area diagrams, which incor-porate chemical and electrochemical thermodynamics simultaneously in a straight-forward manner. These diagrams comprise an extremely useful, yet fundamental,

57M. Bouroushian, Electrochemistry of Metal Chalcogenides, Monographsin Electrochemistry, DOI 10.1007/978-3-642-03967-6_2,C© Springer-Verlag Berlin Heidelberg 2010

58 2 Electrochemistry of the Chalcogens

starting point for the study of electrochemical systems. Certainly, the ability topredict, understand, and ultimately control electrochemical reactions requires alsoknowledge of process kinetics.

Reproduced below are the potential–pH diagrams for the chalcogen–watersystems as originally derived and illustrated by Pourbaix [3], along with someaccompanying data. These will serve as a guide to all subsequent discussion onaqueous systems used for electrochemical preparations of metal chalcogenides. Thereader may take notice of the chalcogen-containing substances that are consideredto be present in the electrolytes and their stability domains. The diagrams repre-sent, almost exclusively, the important “valencies” of the chalcogens, namely –2,+4, and +6 (and not effective oxidation states corresponding, e.g., to polysulfidesand the like). Still, all known substances for each chalcogen, either stable in wateror not, are registered here, in order to serve as a useful basis for the formulae andnomenclature used in the rest of the book. Note that the parts of the Pourbaix dia-grams outside the zero activity lines (log C = 0) regard ideal solutions containing1 M of dissolved chalcogen in the forms considered, and the parts inside these linesregard solutions saturated with solid chalcogen (plus TeO2 in the case of tellurium).The diagrams are valid only in the absence of substances with which the respectivechalcogen can form soluble complexes or insoluble salts.

It is considered useful to include here the potential–pH diagram for some redoxsystems related to oxygen (Fig. 2.1) [4]. Lines 11′ and 33′ correspond to the (a)and (b) dashed lines bounding the stability region of water, as depicted in all thesubsequent Pourbaix diagrams.

Fig. 2.1 Potential–pH diagram for oxygen reactions. (L’ Her [4] Copyright Wiley-VCH VerlagGmbH & Co. KGaA. Reproduced with permission)

2.1 General References 59

2.1.1 Tables of Aqueous Standard and Formal Potentials

Table 2.1 Sulfur reactions

Half-reactionStandard or formal potential(V vs. SHE at 25 ◦C)

S2O2−8 + 2e− → 2SO2−

4 +2.01

S2O2−8 + 2H+ + 2e− → 2 HSO−

4 +2.123

2SO2−4 + 4H+ + 2e− → S2O2−

6 + 2H2O −0.22

SO2−4 + 4H+ + 2e− → H2SO3 + H2O +0.17

SO2−4 + 4H+ + 2e− → SO2 + 2H2O +0.138

SO2−4 + H2O + 2e− → SO2−

3 + 2OH− −0.93HSO−

4 + 7H+ + 6e− → S(s) + 4H2O +0.339

SO2−4 + 8H+ + 6e− → S(s) + 4H2O +0.357

HSO−4 + 9H+ + 8e− → H2S(aq) + 4H2O +0.289

SO2−4 + 10H+ + 8e− → H2S(aq) + 4H2O +0.303

SO2−4 + 9H+ + 8e− → HS− + 4H2O +0.252

SO2−4 + 8H+ + 8e− → S2− + 4H2O +0.149

SO2−4 + 10H+ + 8e− → H2S(g) + 4H2O +0.311

S2O2−6 + 4H+ + 2e− → 2H2SO3 +0.57

S2O2−6 + 2H+ + 2e− → 2HSO−

3 +0.455

S2O2−6 + 2e− → 2SO2−

3 +0.026

3H2SO3 + 2e− → S3O2−6 + 3H2O +0.30

2H2SO3 + H+ + 2e− → HS2O−4 + 2H2O −0.08, −0.056

2SO2−3 +2H2O+2e− → S2O2−

4 +4OH− −1.12

2SO2−3 + 4H+ + 2e− → S2O2−

4 + 2H2O +0.4162HSO−

3 +3H+ +2e− → HS2O−4 +2H2O +0.060

2HSO−3 + 2H+ + 2e− → S2O2−

4 + 2H2O −0.013, −0.009

4H2SO3 + 4H+ + 6e− → S4O2−6 + 6H2O +0.51

4HSO−3 + 8H+ + 6e− → S4O2−

6 + 6H2O +0.5814SO2(g) + 4H+ + 6e− → S4O−2

6 + 2H2O +0.510

2H2SO3 + 2H+ + 4e− → S2O2−3 + 3H2O +0.40

2SO2−3 +3H2O+4e− → S2O2−

3 +6OH− −0.58

2SO2−3 + 6H+ + 4e− → S2O2−

3 + 3H2O +0.705H2SO3 + 2H+ + 2e− → H2SO2 + H2O < +0.4

2HSO−3 + 4H+ + 4e− → S2O2−

3 + 3H2O +0.491

5H2SO3 +8H++10e− → S5O2−6 +9H2O +0.41

H2SO3 + 4H+ + 4e− → S(s) + 3H2O +0.45SO2−

3 + 3H2O + 4e− → S(s) + 6OH− −0.66SO2(g) + 4H+ + 4e− → S(s) + 2H2O +0.451, +0.470

SO2−3 + 6H+ + 6e− → S2− + 3H2O +0.231

S4O2−6 + 2e− → 2S2O2−

3 +0.08, +0.219, −0.10

S4O2−6 + 12H+ + 10e− → 4S(s) + 6H2O +0.416


Table 2.1 (continued)


H2SO2 + 2H+ + 2e− → S(s) + 2H2O >+0.5

S2O2−3 + 6H+ + 4e− → 2S(s) + 3H2O +0.465

S5O2−6 + 12H+ + 10e− → 5S(s) + 6H2O +0.484

SO(g) + 2H+ + 2e− → S(s) + H2O +1.507

5S2O2−3 +30H++24e− → 2S2−

5 +15H2O +0.331

S2O2−3 + 8H+ + 8e− → 2HS− + 3H2O +0.200

S2O2−3 + 6H+ + 8e− → 2S2− + 3H2O −0.006

S2Cl2 + 2e− → 2S(s) + 2Cl− +1.23

5S(s) + 2e− → S2−5 −0.340, −0.315

4S(s) + 2e− → S2−4 −0.33

S(s) + 2H+ + 2e− → H2S(aq) +0.141S(s) + 2H+ + 2e− → H2S(g) +0.171S(s) + H+ + 2e− → HS− −0.065S(s) + H2O + 2e− → HS− + OH− −0.52S(s) + 2e− → S2− A large number of measurements and

calculated values are available. These varyfrom −0.48 to −0.58 V, but most are closerto −0.48 V

4S2−5 + 2e− → 5S2−

4 −0.441

S2−5 + 5H+ + 8e− → 5HS− +0.003

S2−5 + 10H+ + 8e− → 5H2S(g) +0.299

3S2−4 + 2e− → 4S2−

3 −0.478

S2−4 + 2e− → S2− + S2−

3 −0.52

S2−4 + 4H+ + 6e− → 4HS− +0.033

2S2−3 + 2e− → 3S2−

2 −0.506

S2−3 + 2e− → S2− + S2−

2 −0.49

S2−3 + 3H+ + 4e− → 3HS− +0.097

S2−2 + 2e− → 2S2− −0.48, −0.524

S2−2 + 2H+ + 2e− → 2HS− +0.298

Table 2.2 Selenium reactions


HSeO−4 + 3H+ + 2e− → H2SeO3 + H2O +1.090

SeO2−4 + 4H+ + 2e− → H2SeO3 + H2O +1.15

SeO2−4 + 3H+ + 2e− → HSeO−

3 + H2O +1.075

SeO2−4 + 2H+ + 2e− → SeO2−

3 + H2O +0.880

SeO2−4 + 2e− + H2O → SeO2−

3 + 2OH− +0.05




H2SeO3 + 4H+ + 4e− → Se(s) + 3H2O +0.740

HSeO−3 + 5H+ + 4e− → Se(s) + 3H2O +0.778

SeO2−3 + 6H+ + 4e− → Se(s) + 3H2O +0.875

SeO2−3 + 4e− + 3H2O → Se(s) + 6OH− −0.366

Se4+ + 4e− → Se(s) +0.846

H2SeO3 + 6H+ + 6e− → H2Se + 3H2O +0.360

HSeO−3 + 7H+ + 6e− → H2Se + 3H2O +0.386

HSeO−3 + 6H+ + 6e− → HSe− + 3H2O +0.349

SeO2−3 + 7H+ + 6e− → HSe− + 3H2O +0.414

SeO2−3 + 6H+ + 6e− → Se2− + 3H2O +0.276

Se2Cl2 + 2e− → 2Se(s) + 2Cl− +1.1

Se(s) + 2H+ + 2e− → H2Se −0.40

Se(s) + H+ + 2e− → HSe− −0.510

Se(s) + 2e− → Se2− −0.92

Se(s) + 2H+ + 2e− → H2Se(g) −0.369

Table 2.3 Tellurium reactions [TeO2aq(s) refers to hydrated TeO2 crystals]


H2TeO4 + 6H+ + 2e− → Te4++ 4H2O +0.920H2TeO4 + 3H+ + 2e− → HTeO+

2 + 2H2O +0.953H2TeO4 + H+ + 2e− → HTeO−

3 + H2O +0.631HTeO−

4 + 2H+ + 2e− → HTeO−3 + H2O +0.813

HTeO−4 + H+ + 2e− → TeO2−

3 + H2O +0.584

TeO2−4 + 2H+ + 2e− → TeO2−

3 + H2O +0.892H2TeO4 + 2H+ + 2e− → TeO2(s) + 2H2O +1.020H2TeO4 + 2H+ + 2e− → TeO2aq(s) + 2H2O +0.854HTeO−

4 + 3H+ + 2e− → TeO2(s) + 2H2O +1.202HTeO−

4 + 3H+ + 2e− → TeO2aq(s) + 2H2O +1.036

TeO2−4 + 4H+ + 2e− → TeO2(s) + 2H2O +1.509

TeO2−4 + 4H+ + 2e− → TeO2aq(s) + 2H2O +1.343

TeO3(s) + 2H+ + 2e− → TeO2(s) + H2O +1.020TeO3(s) + 2H+ + 2e− → TeO2aq(s) + H2O +0.850Te4+ + 4e− → Te(s) +0.584 (2–3 M HCI), +0.568, +0.556HTeO+

2 + 3H+ + 4e− → Te(s) + 2H2O +0.551H2TeO3 + 4H+ + 4e− → Te(s) + 3H2O +0.589HTeO−

3 + 5H+ + 4e− → Te(s) + 3H2O +0.713

TeO2−3 + 6H+ + 4e− → Te(s) + 3H2O +0.827

TeO2(s) + 4H+ + 4e− → Te(s) + 2H2O +0.521TeO2aq(s) + 4H+ + 4e− → Te(s) + 2H2O +0.604




Te(OH)2−6 + 4e− → Te(s) + 6OH− −0.412

TeCl2−6 + 4e− → Te(s)+ 6Cl− +0.55, +0.630 (2–3 M HCI)

2Te4+ + 10e− → Te2−2 +0.286

2 HTeO+2 + 6H+ + 10e− → Te2−

2 + 4H2O +0.273

2 HTeO−3 + 10H+ + 10e− → Te2−

2 + 6H2O +0.402

2 TeO2−3 + 12H+ + 10e− → Te2−

2 + 6H2O +0.493Te4+ + 2H+ + 6e− → H2Te +0.132HTeO+

2 + 5H+ + 6e− → H2Te + 2H2O +0.121Te2+ + 2e− → Te(s) +0.40Te2S + 2e− → 2Te(s) + S2− −0.90

2Te(s) + 2e− → Te2−2 −0.840, −0.790, −0.74

2Te(s) + 2H+ + 2e− → H2Te2 −0.365 (30 ◦C)

3Te2 + 8e− → 2Te2−2 + 2Te2− −0.92

Te(s) + 2H+ + 2e− → H2Te −0.739, −0.50 (30 ◦C)Te(s) + 2H+ + 2e− → H2Te(g) −0.717Te(s) + 2e− → Te2− −0.913, –1.14

Te2−2 + 4H+ + 2e− → 2H2Te −0.638

Te2−2 + 2H+ + 2e− → 2HTe− −0.795

Te2−2 + 4H+ + 2e− → 2H2Te(g) −0.595

Te2−2 + 2e− → 2Te2− −1.445

2.1.2 Pourbaix Diagram for Sulfur–Water

Solid substances considered: S (sulfur, light yellow, orthorhombic).Dissolved (in aquo) sulfur substances considered:H2S (hydrogen sulfide, colorless), HS– (hydrogen sulfide ion, colorless), S2–

(sulfide ion, colorless), S2−2 (disulfide ion, orange), S2−

3 (trisulfide ion, orange),S2−

4 (tetrasulfide ion, orange), S2−5 (pentasulfide ion, orange), H2S2O3 (thiosul-

furic acid, colorless), HS2O−3 (acid thiosulfate ion, colorless), S2O2−

3 (thiosulfateion, colorless), S5O2−

6 (pentathionate ion, colorless), S4O2−6 (tetrathionate ion, col-

orless), HS2O−4 (acid dithionite ion, colorless), S2O2−

4 (dithionite ion, colorless),S3O2−

6 (trithionate ion, colorless), H2SO3 (sulfurous acid, colorless), HSO−3 (bisul-

fite ion, colorless), SO2−3 (sulfite ion, colorless), S2O2−

6 (dithionate ion, colorless),H2SO4 (sulfuric acid, colorless), HSO−

4 (bisulfate ion, colorless), SO2−4 (sulfate ion,

colorless), S2O2−8 (dipersulfate ion, colorless).

In Fig. 2.2, the potential–pH diagram is represented for the stable equilibriaof the system sulfur–water at 25 ◦C, i.e., equilibria comprising the forms H2S,


Fig. 2.2 Potential–pH diagram for the stable equilibria of the system sulfur–water at 25 ◦C(gaseous hydrogen sulfide is designated in italic letters) (Reproduced from [3], Copyright NACEInternational 2010)

HS–, S2–, S, HSO−4 , SO2−

4 , and S2O2−8 , which contain sulfur only in the oxidation

states –2 and +6 (aside from the solid element); other species, such as thiosulfates,dithionites, sulfites, and polythionates are in “false” equilibrium in aqueous solution.Note also that the persulfates (S2O2−

8 ) are unstable in water, so that if the equilib-ria were attained, only the remaining six forms would be present in solution. Thelimits of the domains of relative predominance of the dissolved substances includedin the Pourbaix diagram (plus solid sulfur) regard the following homogeneous andheterogeneous (solid/liquid, gas/liquid) equilibria, involving redox and non-redoxprocesses:


Limits of the domains of relativepredominance of dissolved substances Redox equilibria

(1′) H2S/HS– (41′) S2O2−8 + 2e– → 2SO2−

4

(2′) HS–/S2– (40′) S2O2−8 + 2H++ 2e– → 2HSO−

4(52′) H2S(g)/H2S(aq) (50′) HSO−

4 + 7H++ 6e– → S(s) + 4H2O(53′) H2S(g)/HS– (51′) SO2−

4 + 8H++ 6e– → S(s) + 4H2O(11′) HSO−

4 /SO2−4 (20′) HSO−

4 + 9H++ 8e– → H2S(aq) + 4H2O(21′) SO2−

4 + 10H++ 8e– → H2S(aq) + 4H2O

(22′) SO2−4 + 9H++ 8e– → HS–+ 4H2O

(23′) SO2−4 + 8H++ 8e– → S2–+ 4H2O

(58′) SO2−4 + 10H++ 8e– → H2S(g) + 4H2O

(42′) S(s) + 2H++ 2e– → H2S(aq)(60) S(s) + 2H++ 2e– → H2S(g)(43′) S(s) + H++ 2e– → HS–

2.1.3 Pourbaix Diagram for Selenium–Water

Solid substances considered: Se (selenium, gray, trigonal).Dissolved (in aquo) selenium substances considered:H2Se (hydrogen selenide, colorless), HSe– (acid telluride ion, colorless), Se2–

(selenide ion, colorless), H2SeO3 (selenous acid, colorless), HSeO−3 (acid selenite

ion, colorless), SeO2−3 (selenite ion, colorless), H2SeO4 (selenic acid, colorless),

HSeO−4 (acid selenate ion, colorless), SeO2−

4 (selenate ion, colorless).


1′ HSe–/H2Se 12′ HSeO−4 + 3H++ 2e– → H2SeO3 + H2O

2′ Se2–/HSe– 13′ SeO2−4 + 4H++ 2e– → H2SeO3 + H2O

3′ HSeO−3 /H2SeO3 14′ SeO2−

4 + 3H++ 2e– → HSeO−3 + H2O

4′ SeO2−3 /HSeO−

3 15′ SeO2−4 + 2H++ 2e– → SeO2−

3 + H2O

6′ SeO2−4 /HSeO−

4 19 H2SeO3 + 4H++ 4e– → Se(s) + 3H2O24 HSe–/H2Se(g) 20 HSeO−

3 + 5H++ 4e– → Se(s) + 3H2O

21 SeO2−3 + 6H++ 4e– → Se(s) + 3H2O

7′ H2SeO3 + 6H++ 6e– → H2Se + 3H2O8′ HSeO−

3 + 7H++ 6e– → H2Se + 3H2O9′ HSeO−

3 + 6H++ 6e– → HSe– + 3H2O

10′ SeO2−3 + 7H++ 6e– → HSe– + 3H2O

11′ SeO2−3 + 6H++ 6e– → Se2– + 3H2O

16 Se(s) + 2H++ 2e– → H2Se17 Se(s) + H++ 2e– → HSe–

18 Se(s) + 2e– → Se2–

22 Se(s) + 2H++ 2e– → H2Se(g)


Fig. 2.3 Potential–pH equilibrium diagram for the system selenium–water, at 25 ◦C (gaseoushydrogen selenide is designated in italic letters) (Reproduced from [3], Copyright NACEInternational 2010)

The potential–pH diagram for the system selenium–water at 25 ◦C is given inFig. 2.3. This diagram was constructed by using the homogeneous and heteroge-neous (solid/liquid, gas/liquid) equilibria listed in the previous page, in which all ofthe above-referred dissolved substances of selenium (as well as solid Se) participate.

2.1.4 Pourbaix Diagram for Tellurium–Water

Solid substances considered:Te (tellurium, brown black, amorphous), TeO2 (tellurous anhydride, white),

TeO2 hydr. (tellurous acid, H2TeO3 or TeO2·H2O, white, amorphous), TeO3 hydr.(orthotelluric acid, H6TeO6 or TeO3·3H2O, colorless, cubic or monoclinic).


Dissolved (in aquo) tellurium substances considered:H2Te (hydrogen telluride, colorless), HTe– (hydrogen telluride ion, colorless),

Te2– (telluride ion, colorless), Te2−2 (ditelluride ion, red), Te4+ (tellurous ion),

HTeO+2 (telluryl ion), HTeO−

3 (acid tellurite ion, colorless), TeO2−3 (tellurite ion,

colorless), H2TeO4 (telluric acid, colorless), HTeO−4 (acid tellurate ion, colorless),

TeO2−4 (tellurate ion, colorless).

The potential–pH diagram for the system tellurium–water at 25 ◦C is given inFig. 2.4. It was constructed by using the following homogeneous and heteroge-neous (solid/liquid, gas/liquid) equilibria, involving redox and non-redox processes,in which all of the above-referred dissolved substances of tellurium, as well as thesolid ones, participate:


1′ H2Te/HTe– 17′ H2TeO4+ 6H++ 2e– → Te4++ 4H2O2′ HTe–/Te2– 18′ H2TeO4+ 3H++ 2e– → HTeO+

2 + 2H2O3′ Te4+/HTeO+

2 19′ H2TeO4+ H++ 2e– → HTeO−3 + H2O

4′ HTeO+2 /HTeO−

3 20′ HTeO−4 + 2H++ 2e– → HTeO−

3 + H2O5′ HTeO−

3 /TeO2−3 21′ HTeO−

4 + H++ 2e– → TeO2−3 + H2O

6′ H2TeO4/HTeO−4 22′ TeO2−

4 + 2H++ 2e– → TeO2−3 + H2O

7′ HTeO−4 /TeO2−

4 38′ H2TeO4 + 2H++ 2e– → TeO2(s) + 2H2O39′ HTeO−

4 + 3H++ 2e– → TeO2(s) + 2H2OTwo dissolved substances in the presence ofTe (solution saturated with elementary Te)

24 TeO3(s) + 2H++ 2e– → TeO2(s) + H2TeO3(s) + 2H++ 2e– → TeO2aq(s) + H2O

34′ Te4++ 4e– → Te(s)8′′ H2Te/Te2−

2 35′ HTeO+2 + 3H++ 4e– → Te(s) + 2H2O

14′′ HTeO+2 /Te2−

2 36′ HTeO−3 + 5H++ 4e– → Te(s) + 3H2O

15′′ HTeO−3 /Te2−

2 37′ TeO2−3 + 6H++ 4e– → Te(s) + 3H2O

16′′ TeO2−3 /Te2−

2 23 TeO2(s) + 4H++ 4e– → Te(s) + 2H2O11′ Te4++ 2H+ + 6e– → H2Te

One solid substance and one dissolvedsubstance

12′ HTeO+2 + 5H++ 6e– → H2Te + 2H2O

33′ 2Te(s) + 2e– → Te2−2

32′ Te(s) + 2H+ + 2e– → H2Te25′ Te4+/TeO2 45′ Te(s) + 2H+ + 2e– → H2Te(g)26′ HTeO+

2 /TeO2 8′ Te2−2 + 4H++ 2e– → 2H2Te

27′ TeO2/HTeO−3 9′ Te2−

2 + 2H++ 2e– → 2HTe–

28′ TeO2/TeO2−3 44′ Te2−

2 + 4H++ 2e– → 2H2Te(g)

One gaseous substance and one dissolvedsubstance

10′ Te2−2 + 2e– → 2Te2–

42′ H2Te(g)/HTe–

2.2 General Discussion 67

Fig. 2.4 Potential–pH equilibrium diagram for the system tellurium–water, at 25 ◦C. (The formof TeO2 considered in this diagram is the anhydrous one, which is more stable than the hydratedform TeO2·H2O. Gaseous hydrogen telluride is designated in italic letters) (Reproduced from [3],Copyright NACE International 2010)

2.2 General Discussion

2.2.1 Sulfur

In the Pourbaix diagram, solid sulfur appears to be stable in a very narrow trian-gular domain, which lies completely within the stability domain of water. Sulfur istherefore stable in the presence of water and in acid solutions free from oxidizingagents. It is unstable, however, in alkaline solutions, in which it tends to dispropor-tionate to give HS–, S2– (and polysulfides), SO2−

4 , and other oxidation products. In


practice, these reactions are slow and take place only in hot, very alkaline media.The sulfides of the alkali metals and the alkaline earth metals are soluble in water,whereas the other sulfides are insoluble but dissolve in varying degrees in acidsolutions.

The various oxidation states of sulfur have been determined by polarography.The electrochemical oxidation of sulfide ions in aqueous solution may lead to theproduction of elementary sulfur, polysulfides, sulfate, dithionate, and thiosulfate,depending on the experimental conditions. Disulfides, sulfoxides, and sulfones aretypical polarographically active organic compounds. It is also found that thiols (mer-captans), thioureas, and thiobarbiturates facilitate oxidation of Hg resulting thus inanodic waves.

The electroreduction behavior of elemental sulfur with formation of polysul-fide anions in classical organic solvents (DMF, DMA, DMSO, CH3CN, etc.) aswell as in liquid ammonia is well documented [5]. Significant progress has beenachieved in the interpretation of this process in dimethylformamide (DMF) [6]. Onthe other hand, electrooxidation of the element to form polysulfur cations is onlyknown in molten salts and in fluorosulfuric acid [7]. The electrochemical dispro-portionation of sulfur has been studied in strongly basic media, e.g., in NaOH+H2Omelts [8, 9]. A large number of investigations regard the electrochemical propertiesof the sulfur–sulfide system in molten mixtures of sulfur and polysulfide [10] ormolten chlorides [11, 12] in connection with the production of alkali sulfur batteries(Chap. 6).

The very low solubility of elemental sulfur in water does not allow the study of itselectrochemical reduction in aqueous media. Thus, the electrochemical behavior ofsulfide (HS−, S2−) and polysulfide ions in water is much less documented than fornon-aqueous solvents, while the relevant investigations bare often speculative con-clusions since the experimental results do not display strong evidence for chemicalspecies involved in the proposed mechanisms. Anyhow, a number of investigationshave been carried out on the electrochemical properties of the aqueous sulfur–sulfidesystem on different substrate electrodes regarding the involved redox behavior, theformation or adsorption of sulfur and sulfur species on the electrodes, and speciationof sulfur species in solution or on the electrode surface. Differences in the methodof formation of sulfur layers, stimulated by diverse interests, have led to apparentlyconflicting observations regarding, for example, the formation or not of polysulfidespecies as intermediates in the oxidation of sulfide ions to sulfur [13–16]. Analogousdispute is found for the electroreduction of sulfur to sulfide ions [17, 18], sincethe understanding of these processes requires the unambiguous identification of thereduced forms of sulfur, i.e., polysulfides.

The surface chemistry of sulfur on metals has been intensively studied as beingextremely important in several technologies, including metallurgy, metal corrosion,and heterogeneous catalysis. For example, alkali sulfides have been widely used fora long time in the flotation of base metal oxide and sulfide minerals [19, 20]; or wellknown are the corrosive effects of sulfide ions on iron based alloys during the pro-cessing of crude oil and natural gas. The adsorption of sulfur on metal surfaces mayhave a remarkable effect on the surface reactivity, notably on the electrochemical


properties. Well documented is the effect of sulfur adsorbent layers on the anodicdissolution rate of metals such as copper, silver, gold, nickel or on the conditions forsilver electrodeposition, the overvoltage for hydrogen evolution on nickel and iron,etc. (see previous references). The blocking effect of sulfur on hydroxyl adsorptionsites of the electrode with subsequent inhibition of the electrochemical oxidationreaction (passivation) is well known, as also is the inhibitory effect of sulfur lay-ers on the electrocatalytic activity of noble metals (especially when the reactants arestrongly adsorbed on the electrode). Interestingly, it has been observed that adsorbedsulfur significantly enhances the electrocatalytic activity of some metals causing anincrease in the reaction rates. The electroreduction of ferric ions on Pt electrodesin acid media has been used as a model reaction for investigation of the effect ofchemisorbed sulfur on the charge transfer across the metal–solution interface. Thiseffect has been interpreted on the basis of the Marcus theory of electron transfer[21, 22].

In the context of environmental chemistry, various polarographic techniques havebeen successfully used to measure elemental sulfur and sulfide (H2S and HS−)concentrations in a variety of media, down to nanomolar concentrations.

Redox reactions of sulfur involving reduction–oxidation of sulfate, sulfide, andsulfur species/clusters are of extreme biological, geochemical, and environmentalimportance. The progress in the electrochemistry of sulfido-clusters is connectedwith better understanding of the processes that occur in living systems with involve-ment of transition metals, since at least seven of those (V, Mo, W, Fe, Ni, Cu,and Zn) are bonded quite often to sulfur (and/or selenium) atoms and are vital forvarious biological functions. For example, the association of sulfur and iron intosimple to more complex molecular assemblies allows a great flexibility of elec-tron transfer relays and catalysis in metalloproteins, which play a major role interminal respiration, photosynthesis, and dinitrogen reduction. Iron–sulfur clustersynthesis in the laboratory and the subsequent study of their electrochemical prop-erties has been stimulated by the occurrence and characterization of such centersin many redox enzymes and provides, in turn, valuable insight into natural reactionmechanisms [23].

2.2.2 Selenium

Selenium is stable in water and in aqueous solutions over the entire pH interval inthe absence of any oxidizing or reducing agent. Selenium can be electrochemicallyreduced to hydrogen selenide or to selenides that are unstable in water and aqueoussolutions. It can be oxidized to selenous acid or selenites and further (electrolyti-cally) to perselenic acid (H2Se2O8). Selenic and selenous acids and their salts arestable in water. The selenides, selenites, and selenates of metals other than the alkalimetals are generally insoluble.

The polarographic behavior of selenium in aqueous solutions has been studiedin great detail, being in fact far simpler than that of sulfur because the sele-nium compounds exhibit electrochemical activity essentially in only two different


valent states, i.e., Se(–II) and Se(+IV) [24, 25]. In polarography, the number andshape of the recorded waves as well as the half-wave potential values are influ-enced considerably by the composition and pH of the solution, mainly throughvariations in the distribution of the possible solution species (H2Se, HSe–, Se2–,SeO2−

3 , HSeO−3 , H2SeO3). Quadrivalent selenium exhibits polarographic waves

in a number of electrolytes including hydrochloric, sulfuric, nitric, and perchloricacids, potassium nitrate, ammonium acetate, ammonia–ammonium chloride, andorthophosphate buffer solutions [26].

As with sulfur, deposition of Se monolayers tends to passivate electrode surfaces.Selenide compound formation involving the electrode material was evident alreadyfrom the early studies. Thus, changes in the polarographic waves of Se(IV) uponstanding were attributed to the formation of “HgSe” compound [27], while otherelectrode materials, such as Ag, Cd, Cu, and In, are all known to react chemicallywith either Se(0) or Se(–II) species. Some electrodes (e.g., Au) promote the UPDof Se(0) via strong metal–Se interactions: Andrews and Johnson [28] reported twowaves for the reduction of Se(IV) at an Au electrode in 0.1 M HClO4; the one at themore positive potential was assigned to a UPD process. The authors argued that theelectrodeposition of Se(IV) produces Se in three distinct states of activity, one byone corresponding to the three anodic stripping peaks observed for large quantitiesof deposited Se. The authors claimed that, in effect, approximately a monolayeris initially deposited which is stabilized by 1D interaction with the Au electrodesurface. The formation of a bulk deposit produces a large activity gradient whichis the driving force for irreversible diffusional transport of Se into the Au electrodeforming Au–Se alloy of unknown stoichiometry. Surface-adsorbed selenium andbulk selenium correspond, respectively, to the other two activity states detected bystripping voltammetry.

The mechanism of Se(IV) reduction has been debated a lot in the last fourdecades. The early polarographic data, as also most recent voltammetric studies,were interpreted in terms of a two-step Se(+IV) → Se(0) → Se(–II) scheme, wherethe reduction of Se(0) takes place at the more negative potentials. However, elec-troreduction of Se(IV), being generally more sluggish than of Te(IV), tends to favora direct six-electron route at relatively positive potentials. In fact, as first recognizedby Skyllas-Kazacos and Miller [29], and emphasized by Wei et al. [30], the 6e– pro-cess generating Se(–II) ions, although not recognized in the majority of the earlierstudies at least in neutral and acidic media, is considered very likely to occur incompetition to the 4e– reduction pathway. Certainly, for a given bath composition,voltammetric or other experiments may reveal which reaction will predominate atany potential.

The observed complexity of the Se(IV) electrochemistry due to adsorptionlayers, formation of surface compounds, coupled chemical reactions, lack of elec-troactivity of reduction products, and other interrelated factors has been discussedextensively. Zuman and Somer [31] have provided a thorough literature-basedreview with almost 170 references on the complex polarographic and voltammet-ric behavior of Se(+IV) (selenous acid), including the acid–base properties, saltand complex formation, chemical reduction and reaction with organic and inorganic


sulfur compounds, polarographic and stripping analysis in the presence of heavymetal ions, as well as the principles and applications of stripping analyses fordetermination of ultratraces of Se(IV). Important information was given also onthe electroanalytical behavior of Se(IV) on gold, silver, platinum, carbon, copper,titanium, and tin oxide electrodes.

Selenium films have been prepared by electrodeposition via both the cathodic andanodic routes. An informative report embodying findings of early practical work hasbeen given by von Hippel and Bloom [32] who used aqueous selenous/selenic acidmixtures as plating baths. These authors pointed out that the main obstacle in elec-troplating selenium lies in the polymorphism of the element, and that unless specialconditions are chosen the material deposits as a non-conductor and interrupts thecurrent flow. They also reported that the electrodeposition of amorphous seleniumpractically ceases in the dark but could be continued under strong illumination.Graham et al. [33] described the deposition of thick amorphous Se deposits fromSeO2 baths containing a surfactant agent. A serious drawback was the short lifeof the bath before defective deposits were obtained. The nature of the cathodematerial was seen to exert a profound influence on the appearance and structureof the deposits. Bright-nickel plating was the best of the substrates examined, whichincluded Pb, Ag, Au, Rh, Cu, dull Ni, bright Cr, black Cr electrodeposits, and Ptsheets. Cattarin et al. [34] studied the Se electrodeposition onto Ti electrodes fromacidic selenite baths, as a function of SeO2 concentration, presence of surfactant,electrode illumination, and temperature. Cyclic voltammetry showed a reductionpeak followed by a “passive” and then a “transpassive” domain at more nega-tive potentials. Growth of the Se film at potentials within the “reduction” domainstopped at a limiting thickness. The film could be grown further in the “transpassive”potential domain.

Trigonal, “metallic” selenium has been investigated as photoelectrode for solarenergy conversion, due to its semiconducting properties. The photoelectrochemistryof the element has been studied in some detail by Gissler [35]. A photodecompo-sition reaction of Se into hydrogen selenide was observed in acidic solutions. Onlyredox couples with a relatively anodic standard potential could prevent dissolutionof Se crystal.

2.2.3 Tellurium

Tellurium is stable in the presence of water and aqueous solutions free from oxi-dizing agents, except in very alkaline solutions, in which it is possible to dissolvewith the simultaneous formation of tellurite, TeO2−

3 , and ditelluride, Te2−2 , ions.

In oxygenated water, tellurium becomes covered with the dioxide TeO2, which isnon-protective if the tellurium is freshly precipitated, but in the compact state itconstitutes a protective film. Tellurium can be cathodically reduced in acid solu-tion to give dissolved and gaseous hydrogen telluride or in alkaline solution to giveTe2−

2 and possibly Te2– ions. It can be anodically oxidized to tellurous anhydrideTeO2 (sparingly soluble at pH around 4–7) or to the telluric oxide (orthotelluric


acid; very soluble), telluric acid, and tellurates. Tellurous ions can be obtained invery acidic solutions. The tellurides, tellurites, and tellurates of metals other thanthe alkali metals are generally insoluble.

On account of its electrochemically active valent states, tellurium exhibits a veryinteresting polarographic behavior. Lingane and Niedrach [25], in their fundamentalstudy on dc polarography and controlled potential coulometry in tellurous acidsolutions, have shown that, except for strongly alkaline medium, the reduction ofTe(+IV) occurs via a four-electron process to tellurium, whereas in sodium hydrox-ide solutions, Te2– results via a 6e– direct reduction. The presence of a sharp andintense maximum on the wave plateau of acidic or neutral pH media was attributedto the Te/Te(–II) couple, but some difficulties in a correct interpretation arose.Regarding telluride solutions, a practically reversible oxidation of Te2– to the ele-mental state was noted in the full range of pH [24]. Shinagawa et al. [36, 37]discussed the Te(IV) reduction mechanism in weakly basic solutions, emphasiz-ing that the two-step reduction to Te(0) and further to Te(–II) at the mercury dropelectrode is complicated by the adsorption of elemental tellurium at the mercurysurface (see also [38–40]). The authors pointed out the influence of UV irradiationon the first reduction step. Panson [41] investigated the polarographic behavior ofthe ditelluride ion prepared by cathodic dissolution of Te in alkaline solution andnoted its disproportionation to Te and Te2–.

Lingane and Niedrach have claimed that the +VI states of tellurium (or sele-nium) are not reduced at the dropping electrode under any of the conditions of theirinvestigation; however, Norton et al. [42] showed that under a variety of conditions,samples of telluric acid prepared by several different procedures do exhibit well-defined (though irreversible) waves, suitable for the analytical determination of theelement. The reduction of Te(+VI) at the dropping electrode was found coulometri-cally to proceed to the –II state (whereas selenate, Se(+VI), was not reduced at thedropping electrode in any of the media reported).

Numerous voltammetric investigations can be found on the electrochemistry oftellurium in aqueous electrolytes and various metals, glassy carbon, or Te electrodes,discussing the mechanism of Te(IV) reduction to Te(0) or the cathodic and anodicstripping of the element (e.g., [43–46]). It is generally accepted that in acid aqueousmedia the Te(+IV) is present in the form of the telluryl ion HTeO+

2 . Experimentaldata for the solubility and diffusivity of HTeO+

2 have been reported by use of theelectrohydrodynamical (EHD) impedance method [47]. In most works, it is assumedthat at low pH, the telluryl ion may be electrochemically reduced to give elementarytellurium via a process that takes place in one step with the transfer of four elec-trons [48]. Further reduction of the formed Te at more negative potentials leads tothe H2Te species and simultaneous hydrogen evolution. In addition, a conpropor-tionation reaction between Te(IV) and Te(–II) appears to give Te(0) deposited onthe electrode. Other authors have proposed that the telluryl ion is electrochemicallyreduced by a direct six-electron transfer to form the H2Te species, which then reactschemically with the telluryl ion to give elementary tellurium on the cathode [49, 50].In any case, using a combination of voltammetry and electrochemical quartz crystalmicrogravimetry (EQCM), Mori et al. [51] argued that the six-electron reduction

References 73

pathway is important and must be considered in the overall mechanistic scheme ofelectroreduction for both Te(IV) and Se(IV).

The possibility that adsorption reactions play an important role in the reduc-tion of telluryl ions has been discussed in several works (Chap. 3; CdTe). Byusing various electrochemical techniques in stationary and non-stationary diffusionregimes, such as voltammetry, chronopotentiometry, and pulsed current electrol-ysis, Montiel-Santillán et al. [52] have shown that the electrochemical reductionof HTeO+

2 in acid sulfate medium (pH 2) on solid tellurium electrodes, gener-ated in situ at 25 ◦C, must be considered as a four-electron process preceded bya slow adsorption step of the telluryl ions; the reduction mechanism was observedto depend on the applied potential, so that at high overpotentials the adsorption stepwas not significant for the overall process.

Data on the electrochemistry of the telluride ion in alkaline media are relativelylimited. Mishra et al. [53] studied the oxidation of Te2– to Te0 at solid electrodes,focusing on the intermediate step(s) of this process, and in particular, the possibilityof detecting ditelluride Te2−

2 via rotating ring disk electrode (RRDE) methodology.Oxidation beyond the elemental state to TeO2−

3 and TeO2−4 was also studied using

cyclic and hydrodynamic voltammetry.Various classes of organotellurium compounds in aprotic solvents have been

studied as to their electrochemistry, whereas more limited are the reports on theelectrochemistry of tellurium and its inorganic compounds in non-aqueous solvents[54, 55].

References

1. Zhdanov I (1975) Sulfur. In Bard AJ (ed) Encyclopedia of electrochemistry of the elements,Marcel Dekker, New York (Vol 4, pp. 273–360); Selenium. Ibid (pp. 361–392); Tellurium.Ibid (pp. 393–443)

2. Latimer WM (1952) The oxidation states of the elements and their potentials in aqueoussolutions. 2nd edn. Prentice Hall, New York

3. Pourbaix M (1974) Atlas of electrochemical equilibria in aqueous solutions. Nationalassociation of corrosion engineers (2nd English Edn.) USA

4. L’Her M (2006) Redox Properties, Electrochemistry of Oxygen. In Bard AJ, Stratmann M(eds) Encyclopedia of Electrochemistry, Vol. 7a: Inorganic Chemistry, Scholz F, Pickett ChJ(eds), Wiley-VCH, Weinheim, p 117

5. Demortier A, Lelieur J-P, Levillain E (2006) Sulfur. In Bard AJ, Stratmann M (eds)Encyclopedia of Electrochemistry, Vol. 7a: Inorganic Chemistry, Scholz F, Pickett ChJ (eds),Wiley-VCH, Weinheim, p 253

6. Levillain E, Gaillard F, Leghie P, Demortier A, Lelieur JP (1997) On the understanding of thereduction of sulfur (S8) in dimethylformamide (DMF). J Electroanal Chem 420: 167–177

7. Fehrmann R, Bjerrum NJ, Poulsen FW (1978) Lower oxidation states of sulfur. 1.Spectrophotometric study of the sulfur-chlorine system in molten sodium chloride-aluminumchloride (37:63 mol%) at 150 ◦C. Inorg Chem 17: 1195–1200

8. Moscardo-Levelut MN, Plichon V (1984) Sulfur chemistry in equimolar NaOH-H2O melt. I.Electrochemical oxidation of sodium sulfide. J Electrochem Soc 131: 1538–1545

9. Claes P, Dewilde Y, Glibert J (1988) Chemical and electrochemical behaviour in moltenalkali hydroxides: Part II. Electrochemistry of chalcogenide ions in the molten NaOH + KOH(49 mol%) eutectic mixture. J Electroanal Chem 250: 327–339


10. Dobson JC, McLarnon FR, Cairns EJ (1986) Voltammetry of sodium polysulfides at metalelectrodes. J Electrochem Soc 133: 2069–2076

11. Warin D, Tomczuk Z, Vissers DR (1983) Electrochemical behavior of Li2S in fused LiCl-KClelectrolytes. J Electrochem Soc 130: 64–70

12. Weaver MJ, Inman D (1975) The sulphur-sulphide electrode in molten salts–I:Chronopotentiometric behaviour in lithium chloride–potassium chloride eutectic. ElectrochimActa 20: 929–936

13. Allen PL, Hickling A (1957) Electrochemistry of sulphur. Part 1 – Overpotential in thedischarge of the sulphide ion. Trans Faraday Soc 53: 1626–1635

14. Briceno A, Chander S (1990) Oxidation of hydrosulphide ions on gold – Part I: A cyclicvoltammetry study. J Appl Electrochem 20: 506–511

15. Wierse DG, Lohrengel MM, Schultze JW (1978) Electrochemical properties of sulfuradsorbed on gold electrodes. J Electroanal Chem 92: 121–131

16. Van Huong CN, Parsons R, Marcus P, Montes S, Oudar J (1981) Electrochemical behaviourof silver and gold single-crystal surfaces covered with a monolayer of adsorbed sulphur.J Electroanal Chem 119: 137–148

17. Hamilton IC, Woods R (1983) An investigation of the deposition and reactions of sulphur ongold electrodes. J Appl Electrochem 13: 783–794

18. Buckley AN, Hamilton IC, Woods R (1987) An investigation of the sulphur(–II)/sulphur(0)system on gold electrodes. J Electroanal Chem 216: 213–227

19. Hamilton IC, Woods R (1981) An investigation of surface oxidation of pyrite and pyrrhotiteby linear potential sweep voltammetry. J Electroanal Chem 118: 327–343

20. Buswell AM, Nicol MJ (2002) Some aspects of the electrochemistry of the flotation ofpyrrhotite. J Appl Electrochem 32: 1321–1329

21. Samec Z, Weber J (1972) Reduction of ferric ion on a rotating platinum electrode of theturbulent type in the presence and absence of adsorbed sulfur. J Electroanal Chem InterfacialElectrochem 38: 115–126

22. Samec Z, Weber J (1973) The influence of chemisorbed sulfur on the kinetic parameters of thereduction of Fe3+ ions on a platinum electrode on the basis of the Marcus theory of electrontransfer. J Electroanal Chem Interfacial Electrochem 44: 229–238

23. Barriére F (2006) Aspects of metallo-sulfur cluster’s electrochemistry. In Bard AJ, StratmannM (eds) Encyclopedia of Electrochemistry, Vol. 7b: Inorganic Chemistry, Scholz F, PickettChJ (eds), Wiley-VCH, Weinheim, p 591

24. Lingane JJ, Niedrach LW (1948) Polarography of selenium and tellurium. I. The –2 States.J Am Chem Soc 70: 4115–4120

25. Lingane JJ, Niedrach LW (1949) Polarography of selenium and tellurium. II. The + 4 States.J Am Chem Soc 71: 196–204

26. Christian GD, Knoblock EC, Purdy WC (1963) Polarography of selenium(IV). Anal Chem35: 1128–1132

27. Christian GD, Knoblock EC, Purdy WC (1965) Use of highly acid supporting electrolytes inpolarography. Observed changes in polarographic waves of selenium(IV) upon standing. AnalChem 37: 425–427

28. Andrews RW, Johnson DC (1975) Voltammetric deposition and stripping of selenium(IV) ata rotating gold-disk electrode in 0.1 M perchloric acid. Anal Chem 47: 294–299

29. Skyllas-Kazacos M, Miller B (1980) Studies in selenous acid reduction and CdSe filmdeposition. J Electrochem Soc 127: 869–873

30. Wei C, Myung N, Rajeshwar K (1994) A combined voltammetry and electrochemical quartzcrystal microgravimetry study of the reduction of aqueous Se(IV) at gold. J Electroanal Chem375: 109–115

31. Zuman P, Somer G (2000) Polarographic and voltammetric behavior of selenous acid and itsuse in analysis. Talanta 51: 645–665

32. von Hippel A, Bloom MC (1950) The electroplating of metallic selenium. J Chem Phys 18:1243–1251

References 75

33. Graham AK, Pinkerton HL, Boyd HJ (1959) Electrodeposition of amorphous selenium.J Electrochem Soc 106: 651–654

34. Cattarin S, Furlanetto F, Musiani MM (1996) Cathodic electrodeposition of Se on Tielectrodes. J Electroanal Chem 415: 123–132

35. Gissler W (1980) Photoelectrochemical investigation on trigonal selenium film electrodes.J Electrochem Soc 127: 1713–1716

36. Shinagawa M, Yano N, Kurosu T (1972) Mechanism and analytical aspects of the polaro-graphic maximum wave of tellurium. Talanta 19: 439–450.

37. Shinagawa M, Sorarnasu N, Mori Y, Okuma T (1977) Studies on the prewave of tellurium andits photo-effect. J Electroanal Chem 75: 809–817

38. Volaire M, Vittori O, Porthault M (1974) Determination du tellure (IV) en milieu acide parpolarographie à tension alternative impulsionnelle et par voltammétrie à balayage linéaire.Anal Chim Acta 71: 185–191

39. Vittori O (1980) Polarographic study of adsorbed tellurium at the hanging and droppingmercury electrodes in 1 M hydrochloric or perchloric acid solutions. Anal Chim Acta 121:315–319

40. Sarala Y, Reddy SJ (1986) Electrochemical reduction of tellurium (IV). J Electroanal Chem214: 179–190

41. Panson AJ (1963) Polarography of the ditelluride ion. J Phys Chem 67: 2177–218042. Norton E, Stoenner RW, Medalia AI (1953) Polarography of tellurium (VI). J Am Chem Soc

75: 1827–183043. Traore M, Moddo R, Vittori O (1988) Electrochemical behaviour of tellurium and silver

telluride at rotating glassy carbon electrode. Electrochim Acta 33: 991–99644. Ngac N, Vittori O, Quarin G (1984) Voltammetric and chronoamperometric studies of tel-

lurium electrodeposition of glassy carbon and gold electrodes. J Electroanal Chem 167:227–235

45. Barbier MJ, Becdelievre AM, Becdelievre J (1978) Electrochemical study of tellurium oxido-reduction in aqueous solutions. J Electroanal Chem 94: 47–57

46. Dergacheva MB, Statsyuk VN, Fogel LA (2001) Electroreduction of tellurium (IV) on solidelectrodes in neutral solutions. Russ J Electrochem 37: 626–628

47. Deslouis C, Maurin G, Pebere N, Tribollet B (1988) Investigation of tellurium electrocrystal-lization by EHD impedance technique. J Appl Electrochem 18: 745–750

48. Yagi I, Nakabayashi S, Uosaki K (1998) In situ optical second harmonic rotational anisotropymeasurements of an Au(111) electrode during electrochemical deposition of tellurium. J PhysChem B 102: 2677–2683

49. Gregory WB, Norton ML, Stickney JL (1990) Thin-layer electrochemical studies of the under-potential deposition of cadmium and tellurium on polycrystalline Au, Pt and Cu electrodes.J Electroanal Chem 293: 85–101

50. Dennison S, Webster S (1992) An investigation into the effect of ionic species on thedeposition of tellurium and the formation of cadmium telluride. J Electroanal Chem 333:287–298

51. Mori E, Baker CK, Reynolds JR, Rajeshwar K (1988) Aqueous electrochemistry of telluriumat glassy carbon and gold: A combined voltammetry-oscillating quartz crystal microgravime-try study. J Electroanal Chem 252: 441–451

52. Montiel-Santillán T, Solorza O, Sánchez H (2002) Electrochemical research on telluriumdeposition from acid sulfate medium. J Solid State Electrochem 6: 433–442

53. Mishra KK, Ham D, Rajeshwar K (1990) Anodic oxidation of telluride ions in aqueous base:A rotating ring-disk electrode study. J Electrochem Soc 137: 3438–3441

54. Liftman Y, Albeck M, Goldsmidt JME, Yarnitsky Ch (1984) The electrochemistry of telluriumin methylene chloride. Electrochim Acta 29: 1673–1678

55. Jeng EGS, Sun IW (1997) Electrochemistry of tellurium(IV) in the basic aluminum chloride-1-methyl-3-ethylimidazolium chloride room temperature molten salt. J Electrochem Soc 144:2369–2374

Documents

Chapter 2 Electrochemistry of the Chalcogens€¦ · Chapter 2 Electrochemistry of ... M. Bouroushian, Electrochemistry of Metal Chalcogenides, Monographs 57 in Electrochemistry,